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Introduction to Chemistry. “The world is full of obvious things. Which nobody by any chance ever observe.” -- Sherlock Holmes. Careful observation is the foundation of chemistry as an experimental science, leading us to question what we have observed – how, what why?

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the world is full of obvious things which nobody by any chance ever observe sherlock holmes
“The world is full of obvious things. Which nobody by any chance ever observe.”-- Sherlock Holmes
  • Careful observation is the foundation of chemistry as an experimental science, leading us to question what we have observed – how, what why?
  • The answers to these questions are sought in experiments, which may be described as observations made under controlled conditions
  • Observation and experimentation – the twin pillars of the scientific method
scientific method
Scientific Method…
  • The scientific way of knowing – often called the scientific method – is sometimes presented as a rigid sequence of events
  • It is not however a rigid path – it is a process of discovery!
  • Discovery begins when we make observations and then try to understand what we have observed by asking key questions and proposing possible answers
  • This process of discovery begins as we design and conduct experiments to test whether our answers to these questions are valid!!
what are the steps of the scientific method
What are the steps of the Scientific Method?
  • Make an observation(s)
  • Propose a hypothesis
  • Design and conduct a controlled experiment
  • Analyze the results
  • Form conclusions
controlled experiment
Controlled Experiment…
  • Experiments should be designed so that the effects of different variables on the behavior of a substance can be studied independently
  • A controlled experiment is when only one variable at a time is changed
  • There are two types of variables in an experiment:
    • Independent variable – the one that is deliberately changed
    • Dependent variable – the thing that changes because of the independent variable
an example
An Example…
  • Students were trying to determine if the amount of a Sodium chloride added to calcium carbonate effected the amount of heat given off
  • Give 3 possible variables for this experiment
  • Give the independent and dependent variables
what is chemistry
What is chemistry?
  • Chemistry is the study and investigation of the structure, composition and property of matter and the changes it undergoes
    • The properties of materials are always related to their structure
    • Hence, structure determines properties
measuring and calculating in science
Measuring and Calculating in Science…
  • Chemistry is a quantitative science because it involves measuring and calculating
      • A measurement must have a number and a scale (called a unit) to be meaningful
  • It can also be a qualitative science because it can involve describing what is happening in a reaction
what makes a measurement
What makes a measurement?
  • In order to make measurements, we must meet three requirements…
  • Know what we are trying to measure
  • Have a standard with which to compare whatever we are measuring
  • Have a method for making comparisons
exact numbers
Exact Numbers…
  • A number that is the result from a definition or an exact count
  • For example – there are 12 apples or π = 3.14
  • All are significant
  • Do not limit the number of sig figs
uncertainty in measurement
Uncertainty in Measurement…
  • A measurement always has some degree of uncertainty
  • The amount of uncertainty depends on the precision of the measuring device
  • In science it is customary to report a measurement by recording all certain digits plus the first uncertain (estimated) digit – these numbers are called the significant figures of a measurement
estimating uncertainty in a measurement
Estimating Uncertainty in a measurement…
  • Remember all measurements are a result of known values and one estimated number
  • When finding the uncertainty in a measurement, we look at the estimated number
  • For example: 0.023
      • The 3 is the estimated number
      • It is in the 1000th place
      • The estimated uncertainty is written +0.001
      • This measurement has a very small uncertainty
rules for counting significant figures
Rules for Counting Significant Figures…
  • Nonzero integers are always significant figures
  • Zeros – there are three classes of zeros
      • Leading zeros precede all the nonzero digits and are not significant figures
      • Captive zeros are between nonzero digits and always count as significant figures
      • Trailing zeros are at the right end of the number and are only significant if the number contains a decimal point
try a few
Try a few…
  • Tell how many sig figs are in each measurement and tell the uncertainty in each measurement
  • 1508 cm
  • 300.0 ft
  • 20.003 g
  • 0.00705 L
bell ringer
Bell Ringer…
  • Tell the number of sig fig
    • 94300
    • 0.000400670
    • 100000.
    • 56.00
  • Tell the uncertainty of each measurement
    • 3.45
    • 6.0
    • 12
    • 4.725
bell ringer1
Bell Ringer…
  • Each of the following are statements from different labs – tell if they are quantitative or qualitative
    • Bubbling
    • Heat given off
    • 23.6 cm wide
    • A strong odor
    • pH of 5.4
    • 273 K
sig figs in mathematical operations
Sig figs in Mathematical Operations…
  • To this point we have learned to count the significant figures in a given number, but we must also consider how uncertainty accumulates as calculations are carried out
rules for sig figs in mathematical operations
Rules for Sig figs in Mathematical Operations…
  • For multiplication or division the number of sig figs in the answer is the same as the number in the least precise measurement used in the calculation
    • 4.56 x 1.4 = 6.38 → 6.4 (correct answer)
  • For addition or subtraction the number of sig figs in the answer has the same number of decimal places as the least precise measurement used in the calculation
    • 12.11 + 18.0 + 1.013 = 31.123 → 31.1 (correct answer)
bell ringer2
Bell Ringer…
  • Carry out the following mathematical operations and give each result with the correct number of significant figures
  • 1.05 x 10-3 / 6.135
  • 21 – 13.8
  • 20 X 23.00
  • 14.75 + 34.25
bell ringer3
Bell Ringer…
  • The actual length of a certain plank is 26.782 cm. Which of the following measurements is the most accurate? Are the measurements precise?
    • 26.5 cm
    • 26.8 cm
    • 26.202 cm
    • 26.98 cm
rules for rounding
Rules for Rounding…
  • In most calculations you will need to round numbers to obtain the correct number of sig figs
  • When rounding, use only the first number to the right of the last significant figure
  • In a series of calculations, carry the extra digits through to the final result, then round
  • If the digit to be removed
      • Is less than 5, then the preceding digit stays the same

1.33 → 1.3

      • Is equal to or greater than 5, the preceding digit is increased by 1

1.36 → 1.4

precision and accuracy
Precision and Accuracy…
  • Two terms often used to describe the reliability of measurements are precision and accuracy
    • Precision – the degree of agreement among several measurements of the same quantity. It also is known as the degree of reproducibility of the measurement
    • Accuracy – the agreement of a particular value with the true value
bell ringer4
Bell Ringer…
  • Decide if the following lab data is accurate or precise or both
    • 13.2 mL, 13.3 mL, 13.1 mL, 13.2 mL
    • The actual value is 13.0 mL
  • Make each of the following have 3 sig figs
    • 34098
    • 0.0003219
    • 7154
    • 76.78
types of errors
Types of errors…
  • There are two types of errors in measurements
    • Random error (indeterminate error) – a measurement that has an equal probability of being high or low. This type of error occurs in estimating the value of the last digit of a measurement
    • Systemic error (determinate error) – occurs in the same directions each time. The measurement is either always too high or too low
in groups
In groups…
  • There are 365 days/year, 24 hours/day, 12 months/year and 60 minutes/hr. Use this data to determine how many minutes are in a month.
  • Now use the following data to calculate the number of minutes in a month: 24 hours/day, 60 minutes/hour, 7 days/week, and 4 weeks/month.
  • Why are these answers different? Which, if any, is more correct and why?
dimensional analysis
Dimensional Analysis…
  • It is often necessary to convert a given result from one system of units to another
  • The best way to do this is by a method called unit factor method OR dimensional analysis
converting from one unit to another
Converting from One Unit to Another…
  • To convert from one unit to another, use the equivalence statement that relates the two units
  • Derive the appropriate unit factor by looking at the direction of the required change (to cancel unwanted units)
  • Multiply the quantity to be converted by the unit factor to give the quantity with desired units
bell ringer5
Bell Ringer…
  • A marathon race is 26 miles and 385 yards.
    • What is the distance in rods
    • What is the distance in meters
    • What is the distance in furlongs?

5.5 yards = 1 rod

40 rods = 1 furlong

8 furlongs = 1 mile

1 meter = 39.37 inches

1 yard = 36 inches

what if there is more than one unit present
What if there is more than one unit present?
  • When more than one unit is present, decide which unit you want to convert first
    • Convert it first
    • Then convert the second unit

***do not get confused!!!


How fast is a car going 35 miles/hour going in yards/second?

1 mile = 1760 yards;1 hour = 60 minutes; 1 minute = 60 seconds

a few problems
A few problems…
  • How many doughnuts can one purchase for $123 if doughnuts cost $3.25/doz?
  • Convert 9.85 L to gal. 1.06 qt = 1.00 L and 4 qt = 1 gal
  • A certain size of nail cost $1.25/lb. What is the cost of 3.25 kg of these nails? 1kg = 2.2 lb
metric system review
Metric System Review…
  • Scientists recognized that long ago a standard system of units had to be adopted if measurements were to be useful
  • The system agreed upon in 1960 was the International System or le Systeme International (SI system)
  • The SI system is based on the metric system and units derived from the metric system
  • Because fundamental units are not always convenient, the SI system employs prefixes to change the size of the unit
derived units
Derived units…
  • Many SI units are combinations of quantities
  • These units are produced by multiplying or dividing standard units
dimensional analysis with metric units
Dimensional Analysis with metric units…

When converting with metric, always use that value of the unit as compared to the base unit

  • Convert 35.4 mm to m
  • Convert 2327.9 cg to kg
  • How many grams are in 53.24 dg?
bell ringer6
Bell Ringer…
  • Why do we use the metric system?
  • Convert 35.4 mm to m
  • Convert 2327.9 cg to kg
  • How many grams are in 53.24 dg?
  • Convert the following:
    • How many inches are in 3.0 meters?
    • A baby weighs 8.5 lbs. How many grams is that?
    • How many gallons of Coke would you drink if you drank entire 2 liter?

Science fiction often uses nautical analogies to describe space travel. If the starship U.S.S. Enterprise is traveling at warp factor 1.71, what is its speed in knots?

Warp 1.71 = 5.00 times the speed of light

The speed of light = 3.00 x 108 m/s

1 knot = 2000 yd/hr

  • The measure of the resistance of an object to a change in its state of motion OR the amount of “stuff” in an object
  • A scale is used to mass an object
mass vs weight
Mass vs. weight…
  • An important point concerning measurements is the relationship between mass and weight
  • Weight is the force gravity exerts on mass, therefore weight varies with the strength of the gravitational field
  • Therefore if you went to the moon your weight would change but not your mass
  • Many times the terms mass and weight are sometimes used interchangeably, although this is incorrect!
  • The derived SI unit of volume is cubic meters (m3)
  • Many times this unit is way too large to be a practical way of expressing volume in a chemistry lab
  • Instead, a smaller unit cubic centimeters (cm3) is used
  • When dealing with the volumes of liquids and gases, the non-SI unit liter(L) is often used
  • Again the liter is often too large so the unit milliliter (mL) is used
  • This means 1 cm3 = 1mL
  • Round the following to 3 sig figs
    • 96747210
    • 91
    • 0.0006589
  • How many sig figs are in each in #1?
  • What is the uncertainty of each measurement?
    • 34.09
    • 6.0222
    • 12
  • What is the difference between precision and accuracy?
  • Convert the following:
    • How many grams are in 548.9 mg?
    • How many feet are in 34.2 m?
    • How many liters are in 2 gallon and 3.4 quarts?
it can be tricky with volume conversions
It can be tricky with volume conversions…
  • How many mL are in 14.65 kL?
  • How many L are in 48.6 cm3?
  • How many dm3 are in 29100 mL?
bell ringer7
Bell Ringer…
  • A piece of metal has the mass of 3.45 kg. What is its mass in g?
  • A container has 2.3 L of gas in it? What is its volume in mL?
  • A container has 750.00 mL of liquid in it. What is its volume in m3?
what is temperature
What is Temperature?
  • A measure of the AVERAGE kinetic energy
  • When looking at the different temperature scales, all are talking about the same height of mercury
temperature conversions
Temperature Conversions…
  • There are three systems used to measure temperature
    • Degrees Fahrenheit (°F)
    • Degrees Celsius (°C)
    • Kelvin (K)
  • Each has a different way of converting between the values
how the equation for f to c was derived
Notice -- 0°C = 32°F and 100°C = 212°F

If we subtract these values then…

100°C = 180°F

* Find the value of 1°C

1°C = (180/100) °F

1°C = 9/5 °F

How the equation for °F to °C was derived…

1. Converting from °C to Kelvin

TC = TK – 273.15

TK = TC + 273.15

  • Converting from °C to °F

TF = TC x 9°F + 32°F


more converting
More Converting…
  • Converting °F to °C

TC = (TF - 32°F)5°C


try these
Try these…

1. Normal body temperature is 98.6°F. Convert this to the Celsius and Kelvin scales.

2. Liquid nitrogen, which is often used as a coolant for low-temperature experiments has a boiling point of 77 K. What is this temperature on the Fahrenheit scale?

  • Ratio of mass to volume
  • D = m/V
  • Unit usually: g/mL or g/cm3
  • Useful for predicting mass
  • Does not depend on the amount of material in a compound
  • A characteristic physical property of a substance
  • It does not depend on the size of the sample because as the sample’s mass increases, its volume increases proportionally
  • Density varies with temperature – V ↑ T↑ D↓
bell ringer8
Bell Ringer…
  • What is the density of an object that has a mass of 14 g and a volume of 2 ml?
  • What is the mass of an object with a volume of 2 cm3 and a density of 1.5 g/mL?
  • What is the volume of an object with the mass of 20 kg and a density of 2.5 g/mL?
bell ringer9
Bell Ringer…
  • An empty container weighs 121.3 g. Filled with carbon tetrachloride (density 1.53 g/cm3) the container weighs 283.2 g. What is the volume of the container?
  • A 55.0 gal drum weighs 75 lbs when empty. What will the total mass be when filled with ethanol?

density of ethanol = 0.789 g/cm3

1 gal = 3.78 L

1 lb = 454 g


In the opening scenes of the movie Raiders of the Lost Ark, Indiana Jones tries to remove a gold idol from a booby-trapped pedestal. He replaces the idol with a bag of sand of approximately the same volume. (density of gold = 19.32 g/mL; density or sand = 2 g/mL)

    • Did he have a reasonable chance of not activating the mass sensitive bobby-trap?
    • In a later scene he and an unscrupulous guide play catch with the idol. Assume the volume of the idol is 1.0 L. If it were solid gold, what mass would the idol have?
      • Is playing catch plausible?
it s all that is matter
It’s all that is Matter…
  • Matter is anything that has mass and takes up space
  • All matter, regardless of form, has some properties in common
      • Volume – the amount of 3-D space an object occupies
      • Mass – a measure of the amount of matter in an object
basic building blocks of matter
Basic Building Blocks of Matter…
  • The most fundamental parts of matter are atoms and molecules, which make up elements and compounds
      • Atom – the smallest unit of an element that has all the properties of that element
      • Element – a pure substance made of only one kind of atom
      • Compound – a substance that is made from the atoms of two or more elements that are chemically bonded
      • Molecule – the smallest unit of an element or compound that retains all the properties of that element or compound
classification of matter
Classification of Matter…
  • Matter is classified according to how it is organized
  • Matter is complex and has different levels of organization:
    • Mixtures
    • Pure substances
  • Most of the matter around us consists of mixtures of pure substances
  • Mixture – consists of materials with variable compositions
  • Two types of mixtures
      • Homogeneous
      • heterogeneous
homogeneous mixture
Homogeneous mixture…
  • Homogeneous mixture – having visibly (to the naked eye) indistinguishable parts
  • Has one phase present
  • It also called a solution
  • There are two parts of a solution

1. the solute – the part that is dissolved

2. the solvent – the part that does the dissolving


  • A homogeneous mixture of metallic elements with one solid phase
alloy problems conversion of total mass to mass of a component
Alloy Problems…Conversion of Total Mass to Mass of a Component
  • Manganese steel is very strong and finds use as railroad rails. It is composed of 86.0% iron, 13.0% manganese, 1.0% carbon. What is the mass of each of the three elements in a 254-kg sample of manganese steel?
one more try
One more try…
  • A sample of brass is composed of 72% copper and the remainder zinc. What mass of brass can be made from 25-kg of zinc?
examples of solutions
Examples of solutions…
  • Air – a gaseous mixture of various gases
  • Brass – a solid mixture of various metals
  • Iced Tea – a liquid mixture of various materials
heterogeneous mixtures
Heterogeneous mixtures…
  • Heterogeneous mixture – having visibly distinguishable parts
  • Can usually be separated into two or more homogeneous mixtures or pure substances
  • Has 2 or more phases present
examples of heterogeneous mixtures
Examples of heterogeneous mixtures…
  • Sand and water
  • Iced tea and ice cubes
  • Pepperoni pizza
pure substances
Pure substances…
  • Mixtures can be separated into pure substances by physical methods
  • Pure substance – contains materials with a constant composition, such as compounds and elements
  • Have a definite composition and definite unchanging properties (both chemical and physical)
classify the following as a mixture or a pure substance
Classify the following as a mixture or a pure substance
  • Salt water
  • Smog
  • Water
  • 10 karat gold
  • Sugar
  • Diamond
  • Coffee
  • Chex mix
  • Chef Salad
  • We use two type of properties to describe people
    • Physical properties – what we look like
    • Emotional/Personality – how we interact with other people

We also use properties to describe matter

    • Physical properties – what the matter looks like
    • Chemical properties – how matter interacts with other matter
physical properties
Physical Properties…
  • A physical property is a characteristic that can be observed or measured without changing the identity of the substance
  • It describes a substance
      • Color
      • State at room temperature
      • Melting point
      • Boiling point
      • Density
      • Specific gravity
physical changes
Physical Changes…
  • A physical change is a change in a substance that does not involve a change in the identity of the substance
  • A change of state is a physical change of a substance from one state to another
      • Solid
      • Liquid
      • Gas
      • Plasma
chemical properties
Chemical Properties…
  • A chemical property relates to a substances ability to undergo changes that transform it into different substances
  • Chemical properties are easiest to see when substances react to form new substances which has different properties than the original substances
chemical change
Chemical change…
  • A chemical change is a change in which one or more substances are converted into different substances
  • It can also be called a chemical reaction
      • Reactants – original substances
      • Products – the new substances formed
  • Chemical changes or reactions form products whose properties differ greatly from the reactants
  • Chemical changes do not affect the total amount of matter (and also the mass) present before and after a reaction
methods for separating mixtures
Methods for Separating Mixtures…
  • There are many different ways to separate mixtures based on their physical properties
    • Distillation – uses the volatility of the components
    • Filtration – used when a mixture consists of a solid and a liquid
    • Chromatography – separates based on speed of movement of the components of the mixture
    • Solubility – uses the amount of solute that dissolves in water at a given temperature
    • Density
    • Melting point
energy transfers
Energy Transfers…
  • Physical and Chemical changes are always accompanied by energy changes
  • One way energy can be transferred is through a temperature difference and is called heat (q)
  • The quantitative measurements in energy changes are expressed in joules (J)
  • Energy is an important concept in chemistry and is a property of all matter
  • The ability to do work
  • All objects possess energy
  • Forms of energy:

1. Chemical energy – released as heat energy

2. Nuclear energy

3. Mechanical energy

4. Electrical Energy

5. Light energy

6. Radiant energy (not a property of objects) – the transfer of

energy through empty space

categories of energy
Categories of Energy…
  • There are two main categories of all energy we see:
    • Potential energy – the energy stored in the object
    • Kinetic energy – the energy due to motion
laws of matter and energy
Laws of Matter and Energy…
  • Law of Conservation of Matter
  • Law of conservation of Energy
  • Law of Conservation of Mass-Energy
  • Mass and energy are interconvertible
    • Matter can change forms
    • Energy can change forms
    • Mass can change into energy
  • However, never can mass or energy be created or destroyed
review questions
Review Questions…
  • What is the Law of Conservation of Energy?
  • What are the two main types of energy?
  • Classify either as a chemical change or a physical change
      • Digestion of a candy bar
      • Melting of ice
      • Formation of clouds
      • Growth of plants
      • Fading of dye in a cloth
a little more about the properties of matter
A little more about the Properties of Matter…
  • Every substance, whether it is an element or compound, has characteristics properties
  • These properties are used to distinguish between substances and to separate them
properties of matter con t
Properties of Matter con’t…
  • Properties can also be characteristics of an entire group as with metals
  • Properties can help identify unknown substances – notice the plural – not just one property can identify a substance
  • Properties are either intensive or extensive
      • Intensive – does not depend on the amount of matter present
      • Extensive – does depend on the amount of matter present
states of matter
States of Matter…
  • Matter exists in three physical states:
  • Gas – also known as vapor

* has no fixed volume or shape

* it takes the shape of the container by either

compression or expansion

  • Liquid – has a definite volume but no specific


  • Solid – is rigid and has a fixed volume and a fixed shape
  • Neither liquids or solids are compressible to any appreciable extent
phase change
Phase Change…
  • The change of matter from one state to another state
  • It is a PHYSICAL change
  • Energy changes ALWAYS accompany a phase change between the three states
changes in state
Changes in state…
  • Melting/freezing – solid liquid
  • Sublimation/deposition – solid gas
  • Vaporization/condensation – liquid gas
a graphical representation
A Graphical Representation…



Energy of the system


some other terms associated with phase changes
Some other terms associated with phase changes…
  • The melting point is the same temperature as the freezing point differing only in the direction from which the phase change is approached
  • Boiling point of a liquid is related to the pressure – if you increase the pressure, you increase the boiling point and visa versa!
compare and contrast the phases of matter
Compare and Contrast the Phases of Matter…





review questions1
Review Questions…
  • Describe the three phases of matter?
  • Classify the following as extensive or intensive – color, mass, length, melting point, ductility
  • Classify the following as a chemical or physical property
      • Reactivity
      • Odor
      • Rusting
      • Stability
      • Expansion
      • Porosity
heat energy
Heat Energy
  • Although there are different forms and types of energy that are important to chemistry, heat energy has the most relevance to chemical changes
  • The most obvious thing about heat energy is that it causes changes in the temperature of matter
  • Just like matter, it too can be quantified – how much heat?
how much heat
How much heat?
  • Heat can be lost or gained by a reaction during a chemical change
    • Exothermic – when a chemical reaction releases heat
    • Endothermic – when a chemical reaction absorbs heat
measuring energy changes
Measuring Energy Changes…
  • Experimentally, energy changes of chemical reactions are measured in a calorimeter
  • To change the temperature of a substance, heat must be added or removed – some substances require little heat to cause a change while others require a great deal of heat
specific heat
Specific Heat…
  • The heat needed to raise the temperature of one gram of a substance by one Celsius degree is called the specific heat (Cp) of a substance
  • Every substance has its own specific heat
  • Water has a specific heat of 4.18 J/g°C while aluminum has a specific heat of 0.900 J/g°C
calculating specific heat
Calculating Specific Heat
  • In calculating specific heat, some items must be understood – the heat lost by one substance is gained by another until equilibrium is reached
  • The equation used:

q = (m)(ΔT) (Cp)

q – energy change

m – mass

ΔT – Tfinal – T initial

Cp – specific heat

some examples
Some examples…

How much heat is lost when a solid aluminum ingot with mass 4100 g cools from 660.0 °C to 25 °C?

q = ?

m = 4100g

ΔT = (660 – 25) = 635°C

Cp = 0.900 J/g°C

q = 4100g 635°C 0.900J = 2.35 x 106 J


try these1
Try these…
  • How much heat is required to raise the temperature of 54.5 g PCl3 from 18.6°C to 79.1°C? Cp = 0.874
  • How much heat is required to raise the temperature of 7.90 x 102 g H2O from 38.4°C to 85.4°C? Cp = 4.18 J/g°C